CHEMICAL
BONDING
Cocaine
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1
Chemistry I – Chapter 8
Chemistry I Honors –
Chapter 12
Chemical Bonding
Problems and questions —
How is a molecule or
polyatomic ion held
together?
Why are atoms distributed at
strange angles?
Why are molecules not flat?
Can we predict the structure?
How is structure related to
chemical and physical
properties?
2
Review of Chemical Bonds
Most bonds are
somewhere in
between ionic
and covalent.
• There are 3 forms of bonding:
• _________—complete transfer
of 1 or more electrons from one
atom to another (one loses, the
other gains) forming oppositely
charged ions that attract one
another
• _________—some valence
electrons shared between
atoms
• _________ – holds atoms of a
metal together
3
The type of bond can usually be calculated by 4
finding the difference in electronegativity of
the two atoms that are going together.
Electronegativity Difference
• If the difference in electronegativities
is between:
– 1.7 to 4.0: Ionic
– 0.3 to 1.7: Polar Covalent
– 0.0 to 0.3: Non-Polar Covalent
Example: NaCl
Na = 0.8, Cl = 3.0
Difference is 2.2, so
this is an ionic bond!
5
Ionic Bonds
6
All those ionic compounds were made
from ionic bonds. We’ve been
through this in great detail already.
Positive cations and the negative
anions are attracted to one another
(remember the Paula Abdul
Principle of Chemistry: Opposites
Attract!)
Therefore, ionic
compounds are usually
between metals and
nonmetals (opposite ends
of the periodic table).
7
Electron
Distribution in
Molecules
G. N. Lewis
1875 - 1946
• Electron distribution is
depicted with Lewis
(electron dot)
structures
• This is how you
decide how many
atoms will bond
covalently!
(In ionic bonds, it
was decided with
charges)
Bond and Lone Pairs
• Valence electrons are distributed
as shared or BOND PAIRS and
unshared or LONE PAIRS.
••
H
Cl
•
•
••
shared or
bond pair
lone pair (LP)
This is called a LEWIS
structure.
8
Bond Formation
A bond can result from an overlap of
atomic orbitals on neighboring atoms.
••
H
+
Cl
••
••
•
•
H
Cl
•
•
••
Overlap of H (1s) and Cl (2p)
Note that each atom has a single,
unpaired electron.
9
Review of Valence Electrons
• Remember from the electron chapter
that valence electrons are the
electrons in the OUTERMOST energy
level… that’s why we did all those
electron configurations!
• B is 1s2 2s2 2p1; so the outer energy
level is 2, and there are 2+1 = 3
electrons in level 2. These are the
valence electrons!
• Br is [Ar] 4s2 3d10 4p5
How many valence electrons are
present?
10
Review of Valence Electrons
Number of valence electrons of a main (A)
group atom = Group number
11
Steps for Building a Dot Structure
Ammonia, NH3
1. Decide on the central atom; never H. Why?
If there is a choice, the central atom is atom of
lowest affinity for electrons. (Most of the time, this is the
least electronegative atom…in advanced chemistry we use a
thing called formal charge to determine the central atom. But
that’s another story!)
Therefore, N is central on this one
2. Add up the number of valence electrons
that can be used.
H = 1 and N = 5
Total = (3 x 1) + 5
= 8 electrons / 4 pairs
12
Building a Dot Structure
3.
Form a single bond
between the central atom and
each surrounding atom (each
bond takes 2 electrons!)
H N H
4.
Remaining electrons form
LONE PAIRS to complete the octet
as needed (or duet in the case of
H).
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8
electrons), while H shares 1 pair.
H
••
H N H
H
13
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Building a Dot Structure
5. Check to make sure there are 8
electrons around each atom
except H. H should only have 2
electrons. This includes SHARED
pairs.
••
H N H
H
6. Also, check the number of electrons in your
drawing with the number of electrons from
step 2. If you have more electrons in the
drawing than in step 2, you must make
double or triple bonds. If you have less
electrons in the drawing than in step 2, you
made a mistake!
15
Carbon Dioxide, CO2
1. Central atom =
2. Valence electrons =
3. Form bonds.
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
This leaves 12 electrons (6 pair).
4. Place lone pairs on outer atoms.
5. Check to see that all atoms have 8 electrons
around it except for H, which can have 2.
Carbon Dioxide, CO2
C 4 eO 6 e- X 2 O’s = 12 eTotal: 16 valence electrons
How many are in the drawing?
6. There are too many electrons in our drawing. We
must form DOUBLE BONDS between C and O.
Instead of sharing only 1 pair, a double bond shares 2
pairs. So one pair is taken away from each atom and
replaced with another bond.
16
Double and
even triple
bonds are
commonly
observed for C,
N, P, O, and S
17
H2CO
SO3
C2F4
18
Now You Try One!
Draw Sulfur Dioxide, SO2
Violations of the Octet Rule
(Honors only)
Usually occurs with B and elements
of higher periods. Common
exceptions are: Be, B, P, S, and Xe.
Be: 4
B: 6
P: 8 OR 10
S: 8, 10, OR 12
BF3
Xe: 8, 10, OR 12
SF4
19
MOLECULAR
GEOMETRY
20
MOLECULAR GEOMETRY
VSEPR
• Valence Shell Electron Pair
Repulsion theory.
• Most important factor in
determining geometry is
relative repulsion between
electron pairs.
Molecule adopts
the shape that
minimizes the
electron pair
repulsions.
21
Some Common Geometries
Linear
Trigonal Planar
Tetrahedral
22
VSEPR charts
• Use the Lewis structure to determine the
geometry of the molecule
• Electron arrangement establishes the bond
angles
• Molecule takes the shape of that portion of
the electron arrangement
• Charts look at the CENTRAL atom for all data!
• Think REGIONS OF ELECTRON DENSITY
rather than bonds (for instance, a double
bond would only be 1 region)
23
24
25
Other VSEPR charts
Structure Determination by VSEPR
Water, H2O
2 bond
pairs
2 lone
pairs
The molecular
geometry is
BENT.
The electron pair
geometry is
TETRAHEDRAL
26
27
Structure Determination by
VSEPR
Ammonia, NH3
The electron pair geometry is tetrahedral.
lone pair of electrons
in tetrahedral position
N
H
H
H
The MOLECULAR GEOMETRY — the
positions of the atoms — is TRIGONAL
PYRAMID.
Bond Polarity
HCl is POLAR because it
has a positive end and a
negative end. (difference
in electronegativity)
+d
-d
••
••
H Cl
••
Cl has a greater share in
bonding electrons than
does H.
Cl has slight negative charge (-d) and H has
slight positive charge (+ d)
28
Bond Polarity
• This is why oil and water will not mix! Oil
is nonpolar, and water is polar.
• The two will repel each other, and so you
can not dissolve one in the other
29
Bond Polarity
• “Like Dissolves Like”
–Polar dissolves Polar
–Nonpolar dissolves
Nonpolar
30
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Diatomic Elements
• These elements do not exist as a single atom;
they always appear as pairs
• When atoms turn into ions, this NO LONGER
HAPPENS!
–
–
–
–
–
–
–
Hydrogen
Nitrogen
Oxygen
Fluorine
Chlorine
Bromine
Iodine
Remember:
BrINClHOF